Notice that in these types of reactions, we denote a certain element or compound
with (s), which stands for "solid". This means that one of the products is not soluble;
it cannot dissolve in solution. When an element or a compound is not soluble, the
element or compound precipitates, meaning it becomes solid and sinks to the bottom
of the solution. Therefore, it is no longer a part of the solution; it exists, but
it is not mixed thoroughly with the solution. Rather, it is a solid sitting within
the solution. Many reactions have at least one element or compound which is aqueous
(aq), meaning that it dissolves in water. However, some solutions also have an element
or compound that is not aqueous (therefore not soluble) in water, which is what
produces a precipitate.

There are rules for telling whether or not an element or compound will precipitate:
these are called Solubility Rules. They are as follows:

1. Salts containing Group I elements are soluble (Li+, Na+,
K+, Cs+, Rb+). Exceptions to this rule are rare.
Salts containing the ammonium ion (NH4+) are also soluble.
2. Salts containing nitrate ion (NO3-) are generally
soluble.3. Salts containing Cl-, Br-, I- are generally
soluble. Important exceptions to this rule are halide salts of Ag+, Pb2+,
and (Hg2)2+. Thus, AgCl, PbBr2, and Hg2Cl2
are all insoluble.
4. Most silver salts are insoluble. AgNO3 and Ag(C2H3O2)
are common soluble salts of silver; virtually anything else is insoluble.
5. Most sulfate salts are soluble. Important exceptions to this rule include
BaSO4, PbSO4, Ag2SO4 and SrSO4.
6. Most hydroxide salts are only slightly soluble. Hydroxide salts of Group
I elements are soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are
slightly soluble. Hydroxide salts of transition metals and Al3+ are insoluble.
Thus, Fe(OH)3, Al(OH)3, Co(OH)2 are not soluble.
7. Most sulfides of transition metals are highly insoluble. Thus, CdS, FeS,
ZnS, Ag2S are all insoluble. Arsenic, antimony, bismuth, and lead sulfides
are also insoluble.
8. Carbonates are frequently insoluble. Group II carbonates (Ca, Sr, and
Ba) are insoluble. Some other insoluble carbonates include FeCO3 and
PbCO3.
9. Chromates are frequently insoluble. Examples: PbCrO4, BaCrO410. Phosphates are frequently insoluble. Examples: Ca3(PO4)2,
Ag3PO411. Fluorides are frequently insoluble. Examples: BaF2, MgF2,
PbF2.

These are known as the solubility rules. Although you may find a slightly different
set of rules if you were to look them up in a text or other resource, they are generally
accepted as listed above, with few, rare variations. You should keep these rules
with you whenever working with chemical equations in order to discern which solutions
will have precipitates and which solutions will not.

Combustion

Form: CxHy + O2 --> CO2 + H2O
+ energy

Example: CH4 + 2O2 --> CO2 + 2H2O + energy

Combustion is also known as burning. It always includes O2 (g) as a reactant-just
like fire needs oxygen to burn, combustion reactions need oxygen to occur. Combustion
reactions also produce heat (released as energy). You can learn more about this
in the Thermodynamics (link) help section. Combustion reactions usually have CxHy
as a reactant and have CO2 and H2O as products. Combustion
reactions almost never go to completion. Practice problems usually assume that combustion
does go into completion; however, when these reactions occur in nature, they reach
equilibrium before they reach completion.

The following reaction is another example of combustion:

2 H2 + O2 --> 2H2O (g) + heat

This reaction is an example of a combustion reaction that does not include carbon.
In this example, the product is gaseous water, or water vapor (also known commonly
as steam).

Acid/Base Reactions

Acid base reactions are neutralization reactions. These reactions occur when an
acid (commonly with a positive hydrogen cation) and a base (commonly with a negative
hydroxide anion) react and bond together to form water. Usually, the other components
react to form a salt, like NaCl.

Bronsted-Lowry Acids and Bases

It should be noted that in the study of
chemistry, there are differing definitions of acids and bases. However,
the most commonly accepted is the Bronsted-Lowry model of acids and bases, which
was composed in 1923 by Johannes Nicolaus Bronsted and Martin Lowry. These two scientists
did not collaborate; rather, they worked independently of each other and came up
with the same theory. Thus, the theory was accepted as true and right. Their model
does not identify acids and bases solely based on the formulation of a salt and
a solvent (water) as we described above. Instead, they define acids as "compounds
having the ability to donate a proton" and bases as "compounds having the ability
to receive a proton." The proton is never donated or received directly from the
nucleus of an atom; rather, it is through the addition or removal of a hydrogen
atom (H+) from a compound.

Here's a simple way of looking at conjugate acids and bases.

A conjugate acid is formed by the addition of a proton in the form of a H+
ion. A common example of a conjugate acid is ammonium, NH4+.
The ammonium ion starts out as NH3, a compound consisting of hydrogen
and nitrogen, and accepts another hydrogen ion, making NH4+.
Now, the ammonium ion is considered to have an "extra" hydrogen ion that it is able
to donate to a compound in need of a proton.

The reaction for this conjugate acid is: NH3 + H+ --> NH4+
(The product NH4+ would be the conjugate acid)

A conjugate base is formed by the removal of a proton in the form of a H+
ion. A common example of a conjugate base is the chloride ion, Cl-. The
chloride ion would start as HCl (hydrochloric acid), a compound consisting of a
hydrogen ion and a chloride ion bonded together. When this compound gives up the
hydrogen ion (H+) it is simply left with the chloride ion, Cl-.
Now, chloride is ready to accept another proton in the form of a hydrogen ion.

The reaction for this conjugate base is: HCl --> H+ + Cl-
(The product Cl- would be the conjugate base)

Redox Reactions (Oxidation-Reduction Reactions)

Form: X --> X+ + e-(oxidation)Y + e- --> Y- (reduction)

Oxidation-reduction reactions, known commonly as redox reactions, are reactions
in which the oxidation state (oxidation number) of atoms changes. More often, electron
transfers between atoms identify these types of reactions. Oxidation is the
loss of electrons, which increases an atom's oxidation state. Reduction is
the gain of electrons, which decreases an atom's oxidation state. On paper, these
are fine definitions; however, in nature, the transfer of electrons may never actually
occur. For example, in reactions with covalent bonds, it is very possible for the
oxidation state to change while never actually experiencing a transfer in electrons.
Therefore, it is better to refer to oxidation as simply an increase in oxidation
state and reduction as a decrease in oxidation state. The following is an example
of a redox reaction:

H2 + F2 --> 2 HF

Redox reactions can be split up into two smaller equations, like this:

H2 ? 2 H+ + 2 e-

and

F2 + 2 e- ? 2 F-

The first smaller equation shows the oxidation of hydrogen. The second smaller equation
shows the reduction of fluorine. Notice that the first reaction splits elemental
hydrogen into hydrogen ions and electrons, making the hydrogen atom more positive.
Also notice that the second reaction combines elemental fluorine with two electrons,
which reduce fluorine by making it more negative. Therefore, when we talk about
this reaction, we can say the following:

Hydrogen is being oxidized.

Fluorine is being reduced.

Fluorine is the oxidizing agent, because it oxidizes hydrogen by accepting two electrons
and is, in turn, reduced.

Hydrogen is the reducing agent, because it reduces fluorine by giving it two electrons
and is, in turn, oxidized.

This may seem really backwards, and from a very basic perspective, it is. The naming
of these does not make good sense, and it can be really tricky to call them the
right things.

Very simply:

The element that is oxidized becomes more positive because it loses its electrons.
The element that is oxidized is known as the reducing agent, because it is helping
the other element be reduced.

The element that is reduced becomes more negative because it gains electrons. The
element that is reduced is known as the oxidizing agent, because it is helping the
other element be oxidized.

Let's think about this. In the above example, hydrogen is giving up two electrons.
When you get rid of electrons, your element becomes more positive, so we can say
that element is oxidized. However, hydrogen is giving the two electrons to fluorine,
which is making fluorine more negative than when we started out. Therefore, fluorine
is reduced (made more negative). Now, since hydrogen is helping to reduce fluorine,
hydrogen is our reducing agent. The element that gives up electrons is always the
reducing agent. Since fluorine is accepting the electrons, it is helping to oxidize
hydrogen. Therefore, fluorine is the oxidizing agent.

Please make a mental note of this now. The reducing agent is oxidized
and the oxidizing agent is reduced. They are absolutely switched around,
and you have to remember this or your completion of redox reactions will seem very
confusing.

Summary of Chemical Reactions

To sum it all up, there are many different types of reactions that occur in nature.
They can be defined as follows:

1. Single Displacement reactions

Form: AB + C --> A + CB

2. Double Displacement reactions

Form: AB + CD --> AD + BC

3. Precipitation reactions

Form: can be single or double displacement reactions (see above) that produce a
solid

4. Combustion

Form: CxHy + O2 --> CO2 + H2O
+ energy

5. Acid/base reactions

Form: acid+ + base- --> salt + water

6. Redox reactions (Oxidation-reduction reactions)

Form: X --> X+ + e- (oxidation)
Y + e- --> Y- (reduction)

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